Cobalt Chemistry

History

The
origin of the name Cobalt
is thought to stem from the German kobold for "evil spirits or goblins",
who were superstitiously thought to cause trouble for miners, since the cobalt
minerals contained arsenic that injured their health and the cobalt ores did not
yield metals when treated using the normal methods. The name could also be
derived from the Greek kobalos for "mine".
Cobalt was discovered in 1735 by the Swedish chemist Georg Brandt.

An excellent site for finding the properties of the elements,
including cobalt is at

The principal ores of Cobalt are cobaltite, [(Co,Fe)AsS], erythrite,
[Co3(AsO4)2.8(H2O)],
glaucodot, [(Co,Fe)AsS], and
skutterudite, [CoAs3].
World production of cobalt has steadily increased in recent years,
almost trebling since 1993. The dominance of African copper-cobalt producers
has been replaced by a more even spread of output between leading
producing countries, with Canada, Norway and more recently Australia,
together with exports from Russia, replacing lost production in the
Democratic Republic of Congo (Zaire). The strongest growth in production of cobalt has
come from Finland, where output grew at over 16% between 1990 and 2002.

Alloys, such as:
Superalloys, for parts in gas turbine aircraft engines.
Corrosion- and wear-resistant alloys. Estimated as about 20% of production in 2003

High-speed steels.

Cemented carbides (also called hard metals) and diamond tools.

Magnets and magnetic recording media.

Catalysts for the petroleum and chemical industries.

electroplating because of its appearance, hardness, and resistance to oxidation.

Drying agents for paints, varnishes, and inks.

Ground coats for porcelain enamels.

Pigments (cobalt blue, known in ancient times, and Cobalt green).

Battery sector (e.g. electrodes) estimated as about 11% of production in 2003.

Steel-belted radial tires.

Cobalt-60 has multiple uses as a gamma ray source:
* It is used in radiotherapy.
* It is used in radiation treatment of foods for sterilization (cold pasteurization).
* It is used in industrial radiography to detect structural flaws in metal parts.

The Cobalt(III) ion forms many stable complexes, which being
inert, are capable of exhibiting various types of isomerism. The
preparation and characterisation of many of these complexes dates
back to the pioneering work of
Werner
and his students.
Coordination theory was developed on the basis of studies
of complexes of the type:

Werner Complexes

[Co(NH3)6]Cl3

yellow

[CoCl(NH3)5]Cl2

red

trans-[CoCl2(NH3)4]Cl

green

cis-[CoCl2(NH3)4]Cl

purple

Another important complex in the history of coordination
chemistry is hexol.
This was the first complex that could be resolved into its optical isomers
that did not contain carbon atoms. Since then, only three or four others have been found.

Recently a structure that Werner apparently misassigned has been determined to be related
to the original hexol although in this case the complex contains 6 Co atoms,
i.e. is hexanuclear. The dark green compound is not resolvable into optical
isomers.

Werner's hexol and "2nd hexol"

A noticeable difference between chromium(III) and cobalt(III)
chemistry is that cobalt complexes are much less susceptible to
hydrolysis, though limited hydrolysis, leading to polynuclear
cobaltammines with bridging OH- groups, is well known.
Other commonly occurring bridging groups are
NH2-, NH2- and
NO2-, which give rise to complexes such as the
bright-blue amide bridged
[(NH3)5Co-NH2-Co(NH3)
5]5+.
In the preparation of cobalt(III) hexaammine salts by the oxidation
in air of cobalt(II) in aqueous ammonia it is possible to
isolate blue
[(NH3)5Co-O2-Co(NH3)
5]4+. This is moderately stable in concentrated
aqueous ammonia and in the solid state but readily decomposes in acid
solutions to Co(II) and O2, while oxidizing agents such as
(S2O8)2- convert it to the
green, paramagnetic
[(NH3)5Co-O2-Co(NH3)
5]5+ (μ300 = 1.7 B.M.).
In the brown compound both cobalt atoms are Co(III) and are joined by a
peroxo group, O22-, this fits with the
observed diamagnetism; in addition the stereochemistry of the
central Co-O-O-Co group is similar to that of
H2O2.
The green compound is less straightforward. Werner thought that it too
involved a peroxo group but in this instance bridging between Co(III) and Co(IV)
atoms.
This could account for the paramagnetism, but EPR evidence shows
that the 2 cobalt atoms are equivalent, and X-ray
evidence shows the central Co-O-O-Co group to be planar with an O-O distance
of 131 pm, which is very close to the 128 pm of the
superoxide, O2-, ion.
A more satisfactory formulation therefore is that of 2 Co(III) atoms
joined by a superoxide bridge.
A range of Co(II) dioxygen complexes are known, some of which are able
to reversibly bind O2 from the air. During WWII, some US aircraft
carriers are reported to have used these complexes as a solid source
for oxy-acetylene welding. By slightly warming the solid complex
the oxygen was released and when cooled again oxygen would be
coordinated again. Unlike an oxygen cylinder the solid would not
explode if hit by a stray bullet!

Co(acac)3 is a green octahedral complex of Co(III). In the case
of Co(II) a comparison can be made to the Ni(II) complexes.
Ni(acac)2 is only found to be monomeric at
temperatures around 200C in non-coordinating solvents such as
n-decane. 6-coordinate monomeric species are formed at room
temperature in solvents such as pyridine, but in the solid state
Ni(acac)2 is a trimer, where each Ni atom is
6-coordinate. Note that Co(acac)2 actually exists as a
tetramer.

[Ni(acac)2]3
[Co(acac)2]4

Cobalt(II) halide complexes with pyridine show structural isomerism.
Addition of pyridine to cobalt(II) chloride in ethanol can produce
blue, purple or pink complexes each having the composition
"CoCl2pyr2". The structures are 4, 5 and 6 coordinate with
either no bridging chlorides or mono- or di- bridged chlorides.

blue-[CoCl2pyr2] CN=4
pink-[CoCl2pyr2] CN=6

See the notes on isomerism
for examples of Co(III) compounds that show linkage and structural isomerism.

The Department of Chemistry, University of the West Indies,
Mona Campus, Kingston 7, Jamaica.Created July 2002. Links checked and/or last
modified 1st September 2017.URL
http://wwwchem.uwimona.edu.jm/courses/cobalt.html